Chapter 5 The Periodic Law

Transcription

1 z Chapter 5 The Periodic Law

2 z Section 5-1 History of the Periodic Table

3 Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals. Repeating patterns are referred to as periodic. created a table in which elements with similar properties were grouped together

4 several empty spaces were left. In 1871 Mendeleev predicted the existence and properties of elements that would fill three of the spaces. By 1886, all three of these elements had been discovered.

5 Properties of Some Elements Predicted By Mendeleev

6 Moseley and the Periodic Law In 1911, the English scientist Henry Moseley discovered that the elements fit into patterns better when they were arranged according to atomic number The Periodic Law -physical and chemical properties of the elements are periodic functions of their atomic numbers.

7 z Periodicity of Atomic Numbers

8 The Modern Periodic Table The Periodic Table is an arrangement of the elements in order of atomic numbers so that elements with similar properties fall in the same column, or group.

9 z Relationship Between Periodicity and Electron Configurations

10 z Section 5-2 Electron Configuration and the Periodic Table

11 Periods and Blocks of the Periodic Table Elements are arranged vertically in the periodic table in groups that share similar chemical properties. Elements are also organized horizontally in rows, or periods. The length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period.

12 The periodic table is divided into four blocks, the s, p, d, and f blocks. The name of each block is determined by the electron sublevel being filled in that block.

13 The period of an element can be determined by the element s electron configuration. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 What is the highest occupied energy level of this element? 4 the element is in period 4 1s 2 2s 2 2p 6 3s 2 3p 5 What is the highest occupied energy level of this element? 3 the element is in period 3

14 The S-block Elements Groups 1 and 2 The elements of Group 1 of the periodic table are known as the alkali metals. In their pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife. ns 1 The elements of Group 2 of the periodic table are called the alkaline-earth metals. Group 2 metals are less reactive than the alkali metals, but are still too reactive to be found in nature in pure form. Used in fireworks ns 2 Click me

15 z Special cases: H and He Hydrogen has an electron configuration of 1s 1, but doesn t share the same properties as the elements of Group 1. It is a nonmetal. It is a gas. It doesn t fit in with any group. BUT it IS very reactive! Click for H balloons Click for Hindenberg news reel

16 z Like the Group 2 elements, helium has an ns2 group configuration. Yet it is part of Group 18. Because its highest occupied energy level is filled by two electrons, helium possesses special chemical stability.

17 a. Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s 2 is located. b. Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (a)?

18 The d Block The d sublevel first appears when n = 3. The 3d sublevel is slightly higher in energy than the 4s sublevel, so these are filled in the order 4s3d. The d-block elements are metals with typical metallic properties and are often referred to as transition elements.

19 Electrons may deviate from the usual sublevel filling in the d- block. S-block electrons may jump up and remain in the d orbitals. Can result in an empty s orbital. Ex: palladium, Pd, [Kr]4d 10 5s 0 The sum of the d and s electrons equals the group number.

20 An element has the electron configuration [Kr]4d 5 5s 1. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group.

21 The p Block consists of all the elements of Groups except helium p-block elements together with the s-block elements are called the main-group elements The properties of the p block vary greatly. p block includes all of the nonmetals except hydrogen and helium. All six of the metalloids are also in the p block. there are eight p-block metals.

22 Group 17 elements are known as the halogens. the most reactive nonmetals react vigorously with most metals to form a types of compounds known as salts The metalloids, or semiconducting elements, are between nonmetals and metals in the p block. The metals of the p block are generally harder and denser than the alkaline-earth metals, but softer and less dense than the d-block metals. Click for F video

23 Without looking at the periodic table, write the outer electron configuration for the Group 14 element in the second period. Then, name the element, and identify it as a metal, nonmetal, or metalloid.

24 The f Block In the periodic table, the f-block elements are wedged between Groups 3 and 4 in the sixth and seventh periods. Their position reflects the fact that they involve the filling of the 4f sublevel. the lanthanides the first row of the f block shiny metals similar in reactivity to the Group 2 alkaline metals

25 the actinides the second row of the f block between Ac, actinium, and Rf, rutherfordium all radioactive

26 Name the block and group in which each of the following elements is located in the periodic table. Then, name each element. Identify each element as a metal, nonmetal, or metalloid. Finally, describe whether each element has high reactivity or low reactivity. a. [Xe]4f 14 5d 9 6s 1 c. [Ne]3s 2 3p 6 b. [Ne]3s 2 3p 5 d. [Xe]4f 6 6s 2

27 z Section 5-3 Electron Configuration and Periodic Properties

28 Atomic Radii Boudaries of an atom are fuzzy, and an atom s radius can vary under different conditions. To compare different atomic radii, they must be measured under specified conditions. Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.

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30 Atomic Radii, continued Atoms are smaller the farther to the right they are found. This is caused by the increasing positive charge of the nucleus, which attracts electrons toward the nucleus. Atoms are larger the farther down in a group they are found. This is caused by the increasing size of the electron cloud around an atom as the number electron sublevels increases.

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32 Try this! Of the elements magnesium, Mg, chlorine, Cl, sodium, Na, and phosphorus, P, which has the largest atomic radius? Explain your answer in terms of trends of the periodic table. Sodium has the largest atomic radius Atomic radii decrease across a period.

33 Ionization Energy An ion is an atom or group of bonded atoms that has a positive or negative charge. Ex: sodium (Na) easily loses an electron to form Na + A + energy à A + + e - Atoms can lose or gain electrons to form ions. Any process that results in the formation of an ion is referred to as ionization.

34 The energy required to remove one electron from a neutral atom of an element is the ionization energy, IE (or first ionization energy, IE 1 ). Generally, ionization energies of the main-group elements increase across each period and decrease down groups

35 The closer an atom is to having 8 valence electrons, the harder it is to remove an e- from that atom. The bigger the atom is, the farther away the valence e- are from the nucleus, which makes it easier to remove an e-.

36 Electrons can continued to be removed from ions Second ionization energy, IE 2, is the amount of energy needed to remove the second electron from an atom Third ionization energy, IE 3, is the amount of energy needed to remove the third electron from an atom The energy to remove another electron is always higher than the previous ionization energy (table pg 147) Electrons feel a stronger pull from the protons because they are closer to the nucleus.

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38 Electron Affinity Electron affinity the energy change that occurs when an electron is acquired by a neutral atom is called the atom s electron affinity Some atoms readily accept electrons, but some atoms do not Atoms that want electrons release energy when they get a new electron Atoms that have to be forced to take another electron require energy to accept the e -

39 Types of Ions Cations positively charged atoms Atoms lose electrons to become cations They have smaller radii because the size of the electron cloud decreases Electrons are pulled closer to nucleus (more p+ than e-) Usually metals Anions negatively charged atoms Atoms gain electrons to become anions They have bigger atomic radii Electrons are not drawn to nucleus as strongly (more e- than p+) Usually nonmetals

40 Valence Electrons Electrons in the outermost energy level (highest) These electrons are the ones that are lost, gained or shared in the formation of chemical compounds They are often in the incompletely formed main-energy levels (s and p)

41 Electronegativity Electronegativity a measure of the ability of an atom in a chemical compound to attract electrons from another atom in a compound Group 1 and 2 are the least electronegative elements Group17, N and O are the most electronegative elements Decreases down a group, increases across a period

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43 Periodic Trends